All of concepts exams so far combined

Radiation that hits earth’s surface

UV, visible, IR (heat)

Greenhouse effect

earth emits longer IR wavelengths 

• Which are absorbed by greenhouse gases → warming atmosphere (essential for habitability) 

• GG’s remit same E IR in random places (space, earth again → global warming) 

Greenhouse gases

gases that absorb IR when photon w/ same E strikes molecule → excited state, Re-emits photn of same E in random directions

To be a greenhouse gas

• molecule must absorb at least 1 wavelength of IR released by earth 

• Have at least 1 polar bond

Bond types in order of inc. En

nonpolar < polar covalent (permanent dipoles) < Ionic

What determines molecule polarity

 bond polarity + molecular geometry (determines if overall molecule is polar/non-polar)

Covalent bonds

not rigid, vibrate

Fluctuating molecular dipoles

• Permanent or temporary

• IR active 

• Creates charge changes 

3  types of fluctuating molecular dipoles

1. Direction of dipole changes 

2. Magnitude of dipole changes 

3. Temporary molecular dipole appears and disappears

Dipoles

Polar → permanent 

Non-polar → temporary, will only have temporary dipoles when vibrating if contains polar bonds

*non-polar bonds → typically IR inactive

IR active

• Vibration → charge fluctuation, E of wavelength absorbed depends on vibration type → bending vs. stretching

• Stronger the bond (“spring”), the more E needed to vibrate it 

• All molecules w/ polar bond have at least 1 vibration which changes molecular dipole

Relationship between transmittance and absorbance

lower transmittance = greater absorbance

CO2 capture

• Capture at source 

• DAC → direct air capture 

*as atmospheric CO2 inc, [ ] of dissolved CO2 inc. in oceans → acidification

C neutral

 no change in amount of CO2 in atmosphere, use fuels w/ less C like MeOH 

• Can be achieved w/ carbon capture

C (-)

decreases amount of CO2 in atmosphere 

• Can be achieved w/ carbon capture

C(+)

inc. CO2 in atmosphere → world currently

Solubility of gases in liquids

Pgas impacts solubility of a gas in a liquid, Cgas= gas solubility - max [ ] of dissolved gas

Henry’s law

Henry’s law constant (KH)

• is different for each particular gas

• Polar gases have > KH 

• CO2 has greater KH than other nonpolar gases ( it reacts w/ water to make carbonic acid which in turn produces more aqueous CO2) 

Partial pressure (Pgas)

the pressure exerted by an individual gas in a mixture 

• Is equal to the pressure that the gas would exert if it were all by itself 

• Cgas is proportional to Pgas 

• When partial pressure of gas above soln inc., solubility of gas also inc.

Dalton’s law of partial pressures

 the P of a mixture of gases is the sum of the partial pressures of component gases

Mole fraction (x)

a unit of [ ] defined as # of moles of a component in a mixture divided by the total # of moles of all components

Solubility impacts of gas in a liquid

1. Temp → inc. T, dec. Cg, inc.  KH

2. Polarity → like dissolves like

Le Chatliers principle

when a system at equilibrium is disturbed (or stressed), it responds by re-establishing equil. to reduce applied stress

Adding aqueous or gas phase reactant/pdt to a system

 system proceeds in direction that consumes part of the added species

Removing aqueous or gas phase reactant or pdt to a system

shifts system in direction that restores part of the removed species 

How oceans capture CO2

ocean acidification

Bronsted-lowry model

involves H+ transfer, good for showing what happens in rxn 

Lewis model

shows transfer of e-, good for showing what happens

Lewis acid

accepts lone pair, cations, neutral molecules with partial + charges

Lewis base

donates lone pair, most often anions or neutral molecules with LP’s

How MOF’s capture CO2

• Physioabsorption → absorption by IMFs

• Chemisorption → formation of chemical bonds 

* indicator → HIn = conj. A, :In- = conj :B 

Infared spectrum for CO2

• Asymmetric stretching = 4.25 µm → more E, smaller wavelength

• Bending = 18 µm → less E, longer wavelength

3 vibrational modes for CO2

Symmetric stretch → IR INACTIVE (no fluctuation in net dipole) 

Asymmetric stretch → IR ACTIVE 

Bending → IR ACTIVE 

Stretch = change in bond length 

Bending = change in bond angle

3 vibrational modes for H2O

3 peaks → therefore ALL IR ACTIVE 

• Asymmetric stretch 

• Symmetric stretch 

• Bend

Why is methane a greenhouse gas

although seen as nonpolar in most situations, it is IR active. Also, any time bonded atoms are different, consider molecule to be non-polar and IR active, its tetrahedral structure allows for C-H vibrate at frequencies that match certain wavelengths of IR radiation 

Wave number (ν̃)

another unit on IR graphs, low ν̃ = low E and therefore long wavelength and vice versa

CO2 → carbonic acid rxn

CO2 → K2CO3 rxn

Why do no other abundant atmospheric gases bind to -OH

Sequestration

storage of CO2 

• Terrestrial → trees 

• Geological → underground 

T and P At surface level

• Pressure = 1 atm 

• Temperature = 298K

Ideal gas law

• PV=nRT

• R=0.08206 (L*atm)/mol*K

Ideal Gases

• Ideal gas law is accurate at reg, T’s and P’s

  • Calc. value agrees with experimental/actual value

  • Becomes inaccurate at  ↑P and ↓T 

• Assumptions based on Kinetic Molecular Theory (KMT) 

• Assumptions: 

• Conditions are not typical or accurate underground w/ higher P

1. Gas particles = point masses = negligible volume

2. G parts move constantly and randomly throughout space they occupy, elastic collisions, collisions w/ container walls is origin of gas pressure (each collision produces a tiny force) 

3. G’s have negligible IMFs (particles are far apart) (we assume that g particles don’t interact)  

4. KEavg of gas molecules ⍺ T (in K), ∴ ↑T, ↑KEavg, m = const, urms must ↑

T  ⍺  V, rises underground

Kinetic Energy

• KE = ½ mu^2 

• KE avg is used to describe gas particles (not all particles move at same speed) 

• Rms = speed of particle w/ KE=to KEavg of all particles in a gas

• KEavg = ½ mu^2rms 

• * urms ↑ w/ ↓MM

• At same T all gases have same KE avg, ↓m, urms must ↑

One mol of ideal gas if P doubles

volume decreases by 50%  ↑P,  ↓V

At high P

• 1) Gas particles are closer together ∴ IMFs are no longer negligible 

• These actions:

  •  reduce the frequency of particle collisions w/ container 

  •  reduce force of particle collisions w/ container 

  •  ∴ less particles moving more slowly 

• PV/RT term is lower at same T for molecules with more IMFs 

• 2) V of gas particles themselves is sig. Fraction of total V occupied by the gas 

  • Gas sample isn’t all “empty space”, particles are not compressible and don't shrink w/ increasing P 

  •  ∴ actual volume is greater than expected

Rigid walls at High P

• Larger volume = particles don’t interact much, behave like an ideal gas 

• Smaller volume = particles interact more, IMF’s cause particles to stick, act. P is lower than expected 

Flexible walls at high P

• at higher external pressure, particles interact more 

• Actual volume is higher than expected

Low T

• G’s deviate from ideal behavior 

• ↓T = ↓KEavg bc KE ⍺ T

• Causes particles to “stick” together 

•  reduce the frequency of particle collisions w/ container 

•  reduce force of particle collisions w/ container 

Rigid walls at lower temp

actual P is lower than expected 

Flexible walls at lower T

 the actual volume is lower than expected

Van der Waals Eqn./Real gas law

• Constants:

  • a = increases as IMFs increase 

  • b = increases as size increases

a(n/vactual)² corrects for the fact that Pactual is lower than that predicted by the ideal gas law due to IMFS

b corrects for the fact that Vactual is greater than predicted by ideal gas law due to the size of the particles

forming a real gas order 

ideal gas → real gas → liquid

Rule for P underground

increases by about 100 atm/km

Rule for T underground

changes very little for  first few hundred meters until about 300m, then rises by 25℃/km 

Pumping CO2 underground

•   ↑P,   ↑T 

• Volumes of gas samples decrease because   ↑P,  ↓V, although T is ⍺ to V, higher P dominates 

• At low enough depths CO2 becomes supercritical fluid

Phase diagrams

• X-axis = T, y=P 

• Depend on 3 factors: 

  • IMF strength 

  • T

  • P = ratio of forces to SA, Pext impacts how close particles are, impacts phase of sub. 

• Heating s→l: KE is high enough to overcome some IMFs so solid melts 

• If KE is high enough it completely overcomes all IMFs and liquids can vaporize or solids sublime 

  • blue = s → l 

  • Red = l → g 

  • G = s → g 

Pressure

• At higher P 

  • Particles are closer together 

  • Substance is more dense 

• At lower P 

  • Particles are farther apart 

  • Substance is less dense 

• For all sub.’s g phase has lowest density and is favored at lower pressure (l and s favored at higher P)

• For all sub.’s boiling and sublimation T’s are greater at higher P (as water boils, more moles of gas above liquid,  ↑n, ↑P above liquid) 

Triple pt

all phases coexist

Supercrit fluid

 has properties of both a liquid and a gas 

• Like a liquid: non-compressible, can dissolve stuff (CO2 = nonpolar solvent) 

• Like gas: ultralow viscosity, can diffuse though tiny cracks in some rocks 

• Super critical water has 10x more energy than hot water or steam

CO2 → Rock

• M2+ w/ SiO4 4- (silicate) 

• Fe2+, Ca2+, Mg2+

• M2+ reacts with CO3 2- (carbonate) to make MCO3

• The MCO3 precipitate will form, storing the CO2 only if the [ ] of M2+ ions and CO32- are sufficiently high enough 

How CO2 forms into rock

• CO2 is acidic and dissolves underground metals releasing M2+ 

• CO32- ions reacts with M2+ ions to make solid MCO3

H2CO3

polyprotic (diprotic) acid, donates 2 H+ ions, so has 2 Ka values, [H3O+] produced in 2nd step of rxn is negligible  

Ka - ionization constant

• The equil. Constant for the process of an acid ionizing in water

• The higher the Ka value, better H+ donor, higher [H3O+], stronger the weak acid 

• pKa = -log Ka

• ↑Ka,  ↓pKa

For polyprotic acids, why does the acid get progressively weaker?: 

As the acid becomes more negative. More difficult to donate H+. HCO3- is ∴ a weaker acid than H2CO3 and has a lower Ka

*exclude liquids and solids when calculating

Kb - base ionization constant

• The equil. Constant for the process of a base ionizing in water 

• ↑Kb, stronger the base, higher [OH-] 

• pKb = -log Kb 

• ↑Kb, ↓pKb

• The acid with the lowest Ka will have the strongest conj. B

*exclude liquids and solids when calculating

Conj. acid base pairs

• Conj. acid base pairs differ in the presence or absence of one proton (H+) 

• Conj. A has one more H+ than conj. B and therefore has a charge that is higher by +1

Relationship between Ka and Kb

• (Ka of Ha)* (Hb of A-) = Kw = 1.0 10^ -14 

• This relationship is true for any conj. Acid-base pair

*This ONLY works for conj. acid-base pairs

5% rule

• When solving for the unknown (x) in an equil. Problem 

• Where a is the initial concentration of weak base or acid 

• You can ignore the x in the denominator if doing so gives a % ionization less than 5%

If the % ionization turns out to be greater than 5%, then you can’t ignore x in the denominator and need to resolve w/ the quadratic equation

Titration: can determine [ ] of H2CO3, HCO3-, and CO32- in water

• Sample can be titrated with a strong base (like NaOH) 

• Or strong acid (like HCl) 

• * has to be strong 

• * Na and Cl are spectator ions 

Which Reaction Arrow to use

• If just a weak acid or weak base in water, use ⇌

• Ex, acetic acid, ammonia 

• If one reactant is “strong” (acid or base), use → 

• Ex. acetic acid w/ NaOH, ammonia with HCl

Acid-base titration

• Titrant: solution containing a known concentration of a reactant. Add w/ a burette 

• Analyte: solution containing a reactant that we want to learn something about  (such as molar mass, pKa, pKb, the primary “form” in a solution at certain pH 

• Equivalence point: a quantity of titrant has been added that is enough to consume all of the analyte, the reaction is complete

Endpoint

the pH at which an indicator changes color, for a titration, an indicator should have an endpoint that coincides w/ the pH at the equivalence point, allows the indicator to correctly indicate that all A or :B has been neutralized 

pH indicator

weak acids (HIn) with a special property, acid form of indicator (HIn) is a different color than its conj. Base form (In-), they work well at determining the pH w/in +/- 1 unit of the unit indicator pKa 

Limited solubility

compounds that form precipitates, ionic compounds w/ carbonate (CO32-) and silicate (SiO4 4-) most often have limited solubility

• A little bit of the compound does dissolve though 

• Ex Mg(OH)2 has limited solubility and forms a “milky” suspension (not a solution), the limited solubility is what makes it a weak base 

• Can be described with a chemical equilibrium (exclude solids) 

• When equil is achieved, solution is saturated w/ ions 

• Equil. Constant for process = Ksp

Solubility Equilibria

sometimes we need to increase the solubility of compounds that have limited solubility 

• Most anions are basic (decent H+ acceptors) 

• Increase in acidity increases the solubility of the compound 

Solubility and precipitate formation

• A certain amount of “precipitates” dissolve 

• If the concentration of dissolved ions are low enough, no precipitate forms 

• Comparison of the reaction quotient (Q) to the compound’s Ksp value can be used to predict if a solid will be produced in a precipitation reaction 

• Ksp has a fixed value at a given temp, the pdt of the 2 ion concentrations at equil. Must have this value regardless of how equilibrium is approached, you can start with just reactants, just products, or a mixture of both

*Ksp excludes solids and liquids

Q and K

• The reaction quotient (Q) has the same form as the equilibrium constant (K) 

• For K, the concentrations must be those at equilibrium 

• For Q, the concentrations can be those at any point in the reaction 

Q vs. K

Q<K → no precipitate, substance will dissolve 

Q>K → precipitate 

Q=K  → at equil, solution is saturated w/ ions an just at the point where if any more is added it will precipitate

Buffers

• Near half equivalence points 

• Solutions which contain appreciable amounts of both a weak acid and its conj. Base ( or a weak base and its conj. Acid )

• Resist changes in pH upon addition of a strong acid or base 

• How they work: 

  • They can neutralize small amounts of added H3O+ and OH- ions 

  • Added H3O+ or OH- are consumed by the buffer and the impact on the pH is minimal 

  • Buffers can be used to maintain a nearly constant pH upon addition of relatively small amounts of strong acid or base 

  • buffers work as they contain existing weak acid or base that soaks up the added acid or base (H3O+/OH-), therefore keeping the pH steady

Common-Ion effect

• Definition: the shift in the position of an equilibrium caused by the presence or addition of an ion taking part in the reaction

• The ionization of any weak base (A-) is suppressed when a significant amount of its conjugate acid (HA) is already present 

• The ionization of any weak acid (HA) is suppressed when a significant amount of its conjugate base (A-) is already present

• This is an application of Le Chatelier’s principle known as the common-ion effect

• The result of this suppression is that the concentration of the species does not change by much

Henderson - hasselbach equation

can be used to calculate the pH of a solution in which the ionization of a weak acid (or base) is suppressed by the presence of a significant amount of its conjugate base (or acid) 

• Is most accurate when the concentrations of conjugate acid and base are similar (within a factor of 10) → significant amounts of conj. Acid-base pairs

Moles and the Henderson-Hasselbach Equation

If HA and A- are present in the same solution, the ratio of their concentrations is also their mole ratio

Buffer Capacity

• Definition: the quantity of acid or base that a pH buffer can neutralize while keeping its pH within a desired range 

• all buffers have a limited capacity of how much H3O+ or OH- they can neutralize before large changes in pH take place 

• Generally, a buffer begins to lose its usefulness if one component is less than 10% of the other

Selection of suitable buffer mixtures

• To select an appropriate buffer, one must know the pH range that they want to maintain 

• From HH equation, we can see that the pH of a buffer depends on 2 factors: 

  • 1) pKa of the acid, this has the greatest influence on the buffer pH 

  • 2) the ratio of [A-]:[HA]

    • To obtain a slightly more acidic buffer, add more weak acid → [A-]:[HA] < 1

    • To obtain a slightly more basic buffer, add more weak base → [A-]:[HA] >1

    • Ideally want [A-]/[HA] o be between 0.1 and 10:

      • log(0.1)=-1

      • log(10)=+1

      • Buffers work best at pH=pKa +/- 1

Solid Types

• Nonmetallic 

• Ionic 

• Metallic

Nonmetallic Solids

• btwn nonmetals or metalloids

2 types: 

• Covalent → covalent bonds (much stronger!)

• Molecular→ held by IMFs 

Ionic Solids

• Btwn ions

• Brittle, crys. Struct, 

• When F app, atoms move ∴ atoms w/ like charge touch, ∴ breaks apart 

Metallic Solids

• btwn metals + metalloids 

• Malleable, crys. Struct. 

• Have weak bonds

• Electrical conductors

Special type of covalent bonding in metals

• M’s share e- w/ mult. Atoms (ex Cu shares 1 e- w/ 12 other atoms) 

• Allows atoms in adj. Layers to slip past each other

• Explains malleability + conductivity

Reg. covalent bonds

share e- pair btwn pair of atoms, σ bond

Allotropes

same element + phys. state, dif. form + properties

Ores

• natural compound/mix.

• pure elements can be extracted from them 

• Mostly composed of ionic solids

Crystalline solid

• part.’s arranged in 3D array 

• Slow cool/precip. → atoms have time to get in pos.

Amorphous solids

• Lack ordered internal structure

• No time for atoms to get into pos. 

Malleability

• To bend w/ F

• M = neutral, no e- repel ∴ lattice don’t break 

• Hcp + fcc = “smoothest” ∴ “slip” easiest

Metallic bond strength factors

1. # val. e-

2. # neighboring atoms 

3. Atom size 

↑size, e- further from nuclei, ↓stability, ↓attract. btwn nuclei ∴ weaker bond

*good conductors, sea of e- 

P.T. trend for atomic r

↓ group, ← period, ↑r

Packing efficiency

•  % of total V of unit cell occupied by spheres 

• 2 closest packed: hcp and ccp 

• Mat.’s of cubic cells tend to form 4-sided crystals 

Packing eff. = ((Vatoms)/(Vunit cell))*100

Square packing

• atoms in 1st layer (a) touch 4 adj. atoms (most eff. w/ 6)

• 2nd layer directly above first = cubic packing (simple cubic) 

• 2nd layer nestled in spaces created by (a layer) and pattern continues abab → bcc